CHEMICAL BONDING

▪ Atoms or ions are held together in molecules or

compounds by chemical bonds.

▪ The type and number of electrons in the outer

electronic shells of atoms or ions are instrumental

in how atoms react with each other to form stable

chemical bonds.

▪ Over the last 150 years scientists developed

several theories to explain why and how elements

combine with each other.

CHEMICAL BONDS

Bonding in Chemistry

• Central theme in chemistry: Why and How atoms

attach together

• This will help us understand how to:

1. Predict the shapes of molecules.

2. Predict properties of substances.

3. Design and build molecules with particular sets of

chemical and physical properties.

Two of the most common substance on our

dining table are salt and granulated sugar

NaCl C12H22O11

The properties of substances are determined in large part

by the chemical bonds that hold their atoms together

CHEMICAL BONDS

Chemical BondsAll chemical reactions involve breaking of some bonds and formation of new ones which yield new products with different properties.

Bonding Theories

• Lewis bond Theory

• Valence Bond Theory

• Molecular Orbital Theory

Gilbert Newton Lewis

Lewis Bonding Theory

• Atoms ONLY come together to produce a more

stable electron configuration.

• Atoms bond together by either transferring or

sharing electrons.

• Many of atoms like to have 8 electrons in their

outer shell.

– Octet rule.

– There are some exceptions to this rule—the key to

remember is to try to get an electron configuration like

a noble gas. Li and Be try to achieve the He electron

arrangement.

Lewis Symbols of Atoms

• Uses symbol of element to represent nucleus and

inner electrons.

• Uses dots around the symbol to represent valence

electrons.

– Puts one electron on each side first, then pair.

• Remember that elements in the same group

have the same number of valence electrons;

therefore, their Lewis dot symbols will look

alike.

Li• Be• •B• •C• •N• •O: :F: :Ne:• •

•

• • • •

•• •• •• ••

••

Valence electrons

Practice to write the Lewis symbol

for Arsenic

• As is in group 15 (5), therefore it has 5

valence electrons.

•

••

••

As

Using Lewis Theory to Predict Chemical

Formulas Compounds

Predict the formula of the compound that forms between

calcium and chlorine.

Draw the Lewis dot symbols

of the elements.

Ca∙∙Cl ∙∙∙

∙ ∙∙ ∙

Transfer all the valance electrons

from the metal to the nonmetal,

adding more of each atom as you

go, until all electrons are lost

from the metal atoms and all

nonmetal atoms have 8 electrons.

Ca∙∙Cl ∙∙∙

∙ ∙∙ ∙

Ca2+

CaCl2

Cl ∙∙∙

∙ ∙∙ ∙

Examples for Lewis representation

of some chemical bonds

F••

••

•• • F•••••••

HF••

••

•• ••

••F•• •• H O

••••••

••

H•H• O••

•

•••

F F

O O••••

O O

O••

•

•

••O••

•

•

••

➢ Total number of valence electrons = 6 + 4 + 6 = 16

➢ Actually 24 electrons needed for completing the octet of each atom

➢ Thus 24 - 16 = 8 electrons are shared.

➢ Since two electrons make a bond, the molecule should have 4

bonds.

➢ The remaining 8 electrons are lone pair electrons.

Information:

Given: CO2

Find: Lewis structure

Solution Map: formula → skeletal →electron distribution → Lewis

Example:

Write the Lewis structure of

CO2.

O C O....

..

..

Practice—Draw Lewis Resonance

Structures for CNO−

(C Is Central with N and O Attached)

C = 4

N = 5

O = 6

(-) = 1

Total = 16 e-

N C O••

••

•• ••N C O•

•••

••

••

Example NO3─

1. Write skeletal structure.

– N is central because it is the

most metallic.

2. Count valence electrons.N = 5

O3 = 3 x 6 = 18

(-) = 1

Total = 24 e-

-

TYPES OF CHEMICAL BONDS

• Ionic bonds

• Covalent bonds

• Metallic bonds

The three possible

types of bonds.

Ionic compounds consist of a cation and an anion

• the formula is always the same as the empirical formula

• the sum of the charges on the cation and anion in each formula unit

must equal zero. Lewis bonding theory is able to explain ionic bonds

very well.

The ionic compound NaCl

Ionic bonding

• Ionic substances are formed when an atom

that loses electrons relatively easily react

with an atom that has a high affinity for

electrons.

ex. metal-nonmetal compound

Chemical Bonds

Ionic bonds are formed by the attraction of

oppositely charged ions.

Ionic Bonds

• Metal to nonmetal.

• Metal loses electrons to form cation.

• Nonmetal gains electrons to form anion.

• The electronegativity between the metal and the nonmetal must be > than 2.

• Ionic bond results from + to − attraction.

– Larger charge = stronger attraction.

– Smaller ion = stronger attraction.

• Lewis theory allows us to predict the correct formulas of ionic compounds.

Ions that pack as spheres in a very regularpattern form crystalline substances .

Formation of an Ionic Solid

• 1. Sublimation of the solid metal

M(s) → M(g) [endothermic]

• 2. Ionization of the metal atoms

M(g) →M+(g) + e- [endothermic]

• 3. Dissociation of the nonmetal

1/2X2(g) → X(g) [endothermic]

Sublimation of Li

Ionization of Li

Dissociation of F2

Electron affinity

of F

Formation

of solid

Lattice Energy Calculations

k: a proportionality constant that depends on the structure of the solid and the electron configuration of the ions

Q1 and Q2: charges on the ions

r: the shortest distance between the centers of cations and anions

)(Energy Lattice 21

r

QQk=

More Gains and Losses

• Can elements lose or gain more than one electron?

• The element magnesium, Mg, in Group 2 can lose two electron and element oxygen in Group 6 can gain two electrons to form stable Nobel gas configurations. The ions can come together to form a crystal structure.

Relative sizes of some ions and

their parent atoms.

Structure of ionic crystals

Different types

of crystals are

formed

depending on

the ionic radii

and the charge

of the ions

involved.

How about the bonds between

atoms that have the same

electronegativity (as in H-H

molecule) or when the

electonegativuty difference is <

1.0 (as in C-H)?

Convalent Bonds—Sharing

• Some atoms are unlikely to lose or gain electrons because the number of electrons in their outer levels makes this difficult.

• Consider the Lewis dot structure of carbon

• The alternative is sharing electrons.

.C... C+4 + 4e-

Covalent Bonds• Often found between two nonmetals.

• Typical of molecular species.

• Atoms bonded together to form molecules.

– Strong attraction.

• Atoms share pairs of electrons to attain octets.

• Molecules generally weakly attracted to each other.

– Observed physical properties of molecular substance due to these attractions.

Covalent Bonding

• Electron are shared by nuclei

The Convalent Bond

• Shared electrons are attracted to the nuclei of both atoms.

• They move back and forth between the outer energy levels of each atom in the covalent bond.

• So, each atom has a stable outer energy level some of the time.

The formation of a bond between

two atoms.

An electron density plot for the H2

molecule shows that the shared electrons

occupy a volume equally distributed over

BOTH Hydrogen atoms.

Electron Density for the H2 molecule

Chemical Bonds

Covalent bonds form when atoms share 2 or

more valence electrons.

Covalent bond strength depends on the

number of electron pairs shared by the

atoms.

single

bond

double

bond

triple

bond< <

Examples of Convalent Bond

• The neutral particle is formed when atoms share electrons is called a molecule

Single Covalent Bonds

• Two atoms share one pair of electrons.

– 2 electrons.

• One atom may have more than one single bond.

F••

••

•• • F•• •••••

HF••

••

•• ••

••F•• •• H O

•• ••••

••

H•H• O••

•

•

••

F F

Double Covalent Bond• Two atoms sharing two pairs of electrons.

– 4 electrons.

• Shorter and stronger than single bond.

O O••••

O••

•

•

••O••

••

••

O O

Chemical Bonds

Bond Polarity

• Bonding between unlike atoms results in unequal

sharing of the electrons.

– One atom pulls the electrons in the bond closer to its

side.

– One end of the bond has larger electron density than the

other.

• The result is bond polarity.

– The end with the larger electron density gets a partial

negative charge and the end that is electron deficient

gets a partial positive charge.H Cl•

•d+ d-

Nonpolar and polar covalent bonds

Probability representations of the

electron sharing in HF.

Trends in electronegativity across

a period and down a group

Nature of bonds and electronegativity

Electronegativity Bonddifference (∆)

∆ > 2 Ionic

0.4 < ∆ < 2 Polar covalent

∆ < 0.4 Covalent

In practice no bond is totally ionic. There will

always be a small amount of electron sharing.

Percent ionic character of chemical bonds as a

function of electronegativity difference

Bond Polarity and Dipole

Moments

Dipole Moment

μ=QR

Q: center of charge of

magnitude

R: distance

Dipole Moment of HF

1D=3.336×10-30 coulomb meter

μ=(1.6×10-19 C)(9.17×10-11 m)=1.47×10-29

=4.4 D for fully ionic

Measured dipole moment=1.83 D

1.83×3.336×10-30=δ(9.17×10-11)

δ=6.66×10-20

Ionic character=1.83/4.4=41.6%

Bond Polarity

0 0.4 2.0 4.0Electronegativity difference

Covalent Ionic

PolarPure

3.0-3.0

= 0.04.0-2.1

= 1.9

3.0-0.9

= 2.1

Polar Molecules and Electric Field

Polarized electron of HCl bond

Molecular Geometry• Molecules are three-dimensional objects.

• We often describe the shape of a molecule

with terms that relate to geometric figures.

• These geometric figures have characteristic

“corners” that indicate the positions of the

surrounding atoms with the central atom in

the center of the figure.

• The geometric figures also have characteristic

angles that we call bond angles.

Valence Shell Electron Pair

Repulsion (VSEPR Model)

• It is used to predict the geometries of molecules formed from nonmetals.

• Postulate: the structure around a given atom is determined principally by minimizing electron pair repulsion.

• The bonding and nonbonding pairs should be positioned as far apart as possible.

Predicting a VSEPR Structure

• Draw Lewis structure.

• Put pairs as far apart as possible.

• Determine positions of atoms from the

way electron pairs are shared.

• Determine the name of molecular

structure from positions of the atoms.

For non-metals compounds, four pairs of

electrons around a given atom prefer prior

to form a tetrahedral geometry to minimize

the electron repulsions.

• Draw the Lewis structure

• Count the pairs of electrons and arrange them to minimize repulsions

• Determine the positions of the atoms

• Name the molecular structure

• Lone pairs require more space than bonding pair.

• The bonding pairs are increasingly squeezed together as the number of lone pairs increases.

• The bonding pair is shared between two nuclei; and the electrons can be close to either nucleus.

• A lone pair is localized on only one nucleus, so both electrons are close to that nucleus only.

Molecular Geometries

Molecular Geometries

Practice drawing these shapes below

Linear TP Tetra TBP Octa

Electron Pairs

COMPOUND is an aggregate of two or more atoms in a

definite arrangement held together by chemical bonds

H2 H2O NH3 CH4

A diatomic molecule contains only two atoms

H2, N2, O2, Br2, HCl, CO

A poly molecule contains more than two atoms

O3, H2O, NH3, CH4

Polarity of Molecules

• In order for a molecule to be polar it must:

1. Have polar bonds.

• Electronegativity difference—theory.

• Bond dipole moments—measured.

2. Have an unsymmetrical shape.

• Vector addition.

• Polarity effects the intermolecular forces

of attraction.

Molecule Polarity

The O—C bond is polar. The bonding

electrons are pulled equally toward both O

ends of the molecule. The net result is a

nonpolar molecule.

Molecule Polarity

The H—O bond is polar. Both sets

of bonding electrons are pulled

toward the O end of the molecule.

The net result is a polar molecule.

Water molecule behaves as if it

had a positive and negative end.

The Covalent Chemical Bond

Bond Energies

• Bond breaking requires energy (endothermic).

• Bond formation releases energy (exothermic).

• DH = SD(bonds broken) - SD(bonds formed)

energy required energy released

Bond Energies

Covalent Bond Energies and

Chemical ReactionsH2+F2→2HF

ΔH=ΣD (bonds broken)-ΣD (bonds formed)

ΔH=DH-H+DF-F-2DH-F=1×432+1×154-

2×565

=-544 kJ

Bond Energy of CH4

Experimental result : 1652 kJ/mol

C(g)+4H(g) →CH4(g) + 1652 kJ/mol

An average C-H bond energy per mole

of C-H bond: 1652/4=413 (kJ/mol)

Metallic Bonding

• The model of metallic bonding can be used to explain the properties of metals.

• The luster, malleability, ductility, and electrical and thermal conductivity are all related to the mobility of the electrons in the solid.

• The strength of the metallic bond varies, depending on the charge and size of the cations, so the melting points and DHfusion of metals vary as well.

IONIC COMPOUNDS vs METALS

BREAKING INORGANIC MATERIAL

SLIP PLANES

ALLOY vs PURE METAL